Mg Desolvation and Intercalation Mechanism at the Mo6S8 Chevrel

Aug 14, 2015 - In this work, we examine the Mg-ion desolvation and intercalation process at the Chevrel phase Mo6S8 cathode surface from first princip...
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Mg Desolvation and Intercalation Mechanism at the Mo6S8 Chevrel Phase Surface Liwen F. Wan,*,†,‡ Brian R. Perdue,†,§ Christopher A. Apblett,†,§ and David Prendergast†,‡ †

Joint Center for Energy Storage Research (JCESR), Lawrence Berkeley National Laboratory, Berkeley, California 94720, United States ‡ The Molecular Foundry, Lawrence Berkeley National Laboratory, Berkeley, California 94720, United States § Sandia National Laboratories, Albuquerque, New Mexico 87123, United States

ABSTRACT: In this work, we examine the Mg-ion desolvation and intercalation process at the Chevrel phase Mo6S8 cathode surface from first principles. It is reported that in electrolytes based on chlorides in tetrahydrofuran (THF), Mg2+ is strongly coordinated by the counterion Cl− and can form singly charged MgCl+ and Mg2Cl3+ species in solution. During cell discharge, intercalation of Mg into the Chevrel phase requires breaking the strong, ionic Mg−Cl bond. Our simulation results indicate that the stripping of Cl− is facilitated by the existence of another cationic species, Mo on the Chevrel phase surface. Once Mg is intercalated, it leaves the counterion, Cl−, on the surface, bound to Mo. It is found that the chlorinated surface presents higher activation barriers to further intercalate Mg. Instead, the chlorinated surface continues to interact with incoming MgCl+ species and form various MgCly surface adsorbates. With certain energy costs, the neutral MgCl2 unit may be released from these surface adsorbates to reopen Mo sites on the surface and permit continuous Mg intercalation. Presuming compatibility of chloride electrolytes with the Mg metal anode, our work implies that finding a compatible cathode material will depend critically on its ability to catalyze Mg−Cl bond breaking. This may explain the success of the Chevrel phase, with its open Mo sites, permitting intercalation of Mg from the halide solutions, whereas higher-voltage transition metal oxides, which typically lack open metal sites, require more weakly coordinating anions in their electrolytes.

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INTRODUCTION Current Li-ion batteries, which have been the primary focus of electrical energy storage development over the past 40 years, are approaching their theoretical limits. Divalent Mg-ion battery technology, with potential high energy density and volumetric capacity (3877 mAh/cm3), may provide alternative solutions to meet the fast-growing global demand for energy storage. The prototype Mg-ion battery system was established in 2000 by Aurbach et al.1 This prototype includes a complex organohaloaluminate/tetrahydrofuran [Mg(AlCl2BuEt)2/THF] solution as the electrolyte with a Mg metal anode and Mo6S8 cathode to form a rechargeable electrochemical cell. Despite the remarkable success of this prototype system, its practical application is limited by the low operating voltage, the low capacity, the narrow electrochemical window of the electrolyte, and the slow ion dissolution and deposition process across the electrolyte−electrode interfaces.2 In recent years, significant © 2015 American Chemical Society

efforts have been devoted to the development of new cathode materials with an increased Mg intercalation voltage.3−5 However, many of these systems still suffer from sluggish ion diffusion within the bulk materials, not to mention any impedance caused by surface-related interactions. The success of the prototype Mg-ion battery relies on the use of the Chevrel phase (CP), Mo6S8, as the intercalation cathode. The unit cell of CP comprises an Mo6 octahedral cluster that is embedded in a slightly distorted cube whose corners are defined by chalcogen atoms. Between the unit cells, there are open channels that allow for relatively fast and reversible ion diffusion. The theoretical capacity for Mg intercalation is achieved with two Mg ions intercalated per CP formula unit: Received: May 20, 2015 Revised: August 13, 2015 Published: August 14, 2015 5932

DOI: 10.1021/acs.chemmater.5b01907 Chem. Mater. 2015, 27, 5932−5940

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surfaces, all atoms at the bottom layer are fixed and only the atoms at the top layer are allowed to move. The effective screening medium (ESM) model,22 as implemented in QuantumESPRESSO, is used to treat the electrified CP (100) surface and to examine the behavior of the absorbed solvent molecule in the presence of surface electric fields. The electrochemical environment is simulated by placing our CP slab, with the (100) surface exposed on both sides, within asymmetric analytic boundary conditions defined (within the ESM), vacuum (ε = 1) on one side and metal (ε = ∞) on the other. When electrons are added or removed from the slab, the same amount of counter charge is added to the metal (the imaginary electrode). A detailed description of the methodology can be found in ref 22. Within this setup, one THF molecule is placed on a 2 × 2 CP (100) surface (∼13 Å × 13 Å). Relaxations to minima of total energy are performed for various charge states of the slab+THF system (controlled by adjusting the total number of electrons in the calculation), and the resulting THF binding energy is calculated. Note that, because there is no electrolyte with finite ε included, the calculated THF binding energy will likely be overestimated, yet its response to induced surface charge should be qualitatively correct. Upon simulation of the activation barrier to intercalate Mg from various MgxCly+ species, the nudged elastic band (NEB) method is used within the ESM model to accurately describe the electrostatic boundary conditions. Here, a single MgxCly+ unit is placed on a 2 × 2 × 2 CP substrate, and above the substrate, a vacuum of ∼20 Å is added to avoid any fictitious surface interactions across the boundary. In addition, one extra electron is provided to the substrate to mimic the electron-rich environment of the cathode material during cell discharge. Because both MgCl+ and Mg2Cl3+ are singly charged and the substrate is electron abundant, our entire simulation supercells are always neutral. The advantage of using the ESM approach instead of the conventional slab model, in this case, is that explicit dipole corrections (and beyond) are included. Experimental Setup. For full cell tests, a slurry preparation of Mo6S8, carbon, and polyvinylidene difluoride (PVDF) in excess 1methyl-2-pyrrolidinone (NMP) [90:5:5 (wt %)] was used to coat the cathode foils. Once the slurry was mixed to the proper consistency, it was applied to a brushed Ni foil with a doctor blade to a wet thickness of 4 mil. Coatings were dried overnight and then placed in a vacuum oven at 120 °C for 24 h. The coated materials were then punched into 5 /8 in. disks to be assembled into 2032 coin cells. Note that to avoid any potential corrosion to the current collector, we use Ni instead of Al foil. PhMgCl/THF and AlCl3 (99.99%, anhydrous) were purchased from Sigma-Aldrich and used as received. Equal volume solutions of 0.8 M PhMgCl and 0.4 M AlCl3 were prepared by the addition of dry THF at room temperature. After each solution was well mixed, the colorless AlCl3 solution was slowly added to the dark brown PhMgCl solution, yielding a clear, light-brown solution with concentrations of 0.4 M PhMgCl and 0.2 M AlCl3. This solution was mixed by being gently swirled and designated as 0.4 M APC. Coin cells were fabricated using 2032 hardware (Pred Materials). The coin cells consisted of a 5/8 in. Mo6S8 cathode disk, a 50 μm separator (Tonen), and a 5/8 in. abraded Mg metal anode (Leico Industries, 99.9%) followed by a stainless steel backing plate and a wave spring and ∼210 μL of electrolyte and sealed in a coin cell crimper. Fabricated coin cells were then cycled on battery testers (MACCOR), at a rate of 0.3 C, in a temperature-controlled environment at 60 °C. In this work, Crate is defined as the current required to charge the mass of active material to the theoretical maximal capacity of 122 mAh/g in 1 h. Note that our cell design differs from previous work23 in one key respect: we used a standard batterytype separator (Tonen) with a significantly smaller pore structure (20−30 nm, vs 1−3 μm for the glass mat). These separators are also much thinner than a glass mat separator (50 μm vs 250 μm).

one at a lower-energy “inner site” and the other at an “outer site”. Pristine CP also exhibits a metallic electronic ground state, providing fast and localized electronic screening around intercalated Mg2+ ions.6 However, despite our current knowledge of the bulk properties of CP, we still lack a clear understanding of how its surface interacts with various electrolyte components. Previous work implies that Mg ions are strongly coordinated in bulk electrolytes because of the strong electrostatic interactions between the small Mg2+ ion and surrounding species.7 In aqueous electrolytes, excluding high concentrations, both Mg2+ and Cl− ions can be fully solvated by water molecules, which provide strong coordination via oxygen lone pairs and hydrogen bonds, respectively. Organic solvent molecules, such as THF, lack significantly localized (i.e., atomic scale) nucleophilic sites to readily coordinate the electronegative Cl−. The nonaqueous halo-aluminate electrolytes, such as the dichloro complex (DCC),1 the all-phenyl complex (APC),8 and the magnesium aluminum chloride complex (MACC)9 solutions, present coordination complexes comprising both halide anions and organic solvent (THF) molecules. Predominantly, Cl− ions form strong ionic bonds with either Mg2+ or Al3+. It is expected that the singly charged MgxCly+/ nTHF species are the active electrochemical components at the anode− and cathode−electrolyte interfaces, with recrystallization of working electrolytes favoring the dimer, Mg2Cl3+, as the dominant charged species.2 Using density functional theory calculations of an isolated cluster model, we estimate that the Mg−Cl bond strength can easily exceed 3 eV even for neutral clusters. During discharge, as the strongly coordinated Mg ion approaches the electrolyte−cathode interface, how does Mg2+ lose its ionically bound counterion Cl− (at reasonable temperatures) and intercalate into the cathode material? Unlike the charging condition, where external bias is applied to strip and desolvate Mg from these heavily solvated MgxCly/THF complexes and plate it onto the Mg anode, the stripping of Cl− at the electrolyte−cathode interface should be a spontaneous process during discharge. Our calculations indicate, at least for CP, that the intrinsic surface properties of the cathode material play an important role in weakening the bond between Mg2+ and Cl−. In what follows, we examine the surface properties of Mo6S8 and propose possible Mg intercalation mechanisms through the Mo6S8 surface. The proposed chemical reactions that occur at the CP surface are then rationalized by reference to CP cathode surface characterization from experimentally cycled full cells.

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METHODS

Simulation Details. The electronic structure calculations are based on density functional theory (DFT) within the plane-wave, pseudopotential framework under periodic boundary conditions.10,11 We use the Perdew−Burke−Ernzerhof (PBE) form of the generalizedgradient approximation (GGA) to calculate the exchange-correlation energy.12 The electron−ion interactions are captured using pseudopotentials of either the projector-augmented-wave (PAW)13,14 form (in VASP)15,16 or the ultrasoft type17,18 (in QuantumESPRESSO19). Our calculations are sufficiently numerically converged with respect to plane-wave kinetic energy cutoff and Brillouin zone (k-point) sampling. The structural optimizations are performed in VASP, and the structures are considered to be optimized when the residual forces on each atom are 0.1 e/Å2) will be built up upon the arrival of MgCl+. According to the linear relation found in Figure 4, this localized surface charge of 0.1 e/Å2 is expected to be sufficient to remove surface-adsorbed THF molecules in the vicinity. In the current simulation setup, we have ignored the additional dielectric screening of the solvent adjacent to the CP surface. In reality, the CP surface will be screened by the electrolyte and the resulting charge distribution on the CP surface will be different. For the same reason, the THF adsorption enthalpy will be reduced compared to that under our simplified vacuum condition. Nevertheless, upon the arrival of charged MgxCly species, extra charge will be induced on the CP surface that can help to displace the originally chemisorbed THF molecules and permit Mg-ion intercalation. Mg Intercalation through a Clean Mo6S8 (100) Surface. We imagine that charged Mg complexes can appear in the form of either monomers or dimers in a THF solution. During cell discharge, positively charged MgCl+ and Mg2Cl3+ are attracted by the electron-rich CP surface and initiate Mg intercalation. Using a simple cluster model, it is found that Mg2+−solvent interactions are much weaker than Mg2+−anion interactions. For instance, in the neutral MgCl2/2THF cluster, the dissociation energies of Mg−Cl and Mg−THF species are ∼3 and 0.8 eV, respectively. Therefore, in the following tests, we omit the effects of the solvent on Mg intercalation, because the solvent molecules may already have been displaced or desolvated at this stage, permitting direct contact between Mg

Figure 3. ESM model for charged surface calculation: (top) system setup, (middle) induced charge density integrated into the plane perpendicular to the surface, and (bottom) electrostatic potential for variations in surface charge.

media with one representing vacuum (εleft = 1) and the other representing an ideal metal (εright = ∞). By reformulating the solutions to Poisson’s equation, which defines the electrostatic (Hartree) potential in the Kohn−Sham equations, to include the analytic Green’s functions for these asymmetric boundary conditions, we determine the electronic structure of the system for varying surface charges. This is achieved by directly adding or removing electrons from the slab, while maintaining charge neutrality of the entire simulation cell by simultaneously adjusting the compensating charge of the imaginary electrode in the ESM model. From the middle frame of Figure 3, the magnitude of the induced surface charge density increases as more charge is added or removed from the supercell. On the basis of this ESM approach, we can evaluate the adsorption energy of a THF molecule on the CP (100) surface using the following equation: Ead = Etot − ECP − E THF + qE F + ΔV

(1)

where Etot is the total energy of the supercell, ECP is the energy for a clean (100) surface, and ETHF is the energy of an isolated (gas phase) THF molecule. As in calculations of the formation 5935

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Chemistry of Materials and the CP surface. Again, unlike in aqueous solutions where the counterion can also be readily solvated by the water molecule, in organic solvents, the counterion Cl− is not readily solvated. During the desolvation process of Mg2+, we expect the MgxCly+ units will first lose their coordinating organic solvent molecules, and the most important question is how to break the much stronger cation−anion bonds at the cathode surface to allow Mg to intercalate as an isolated unit. Starting from the simplest case, we examine Mg intercalation behavior for MgCl+. As shown in Figure 5, one MgCl+ unit is

Figure 6. Relative formation energy of Mg2Cl3+ on the CP (100) surface. The reference structure (c) denotes the lowest-energy configuration.

Figure 5. Calculated activation energy to intercalate Mg from a MgCl+ unit.

placed on a 2 × 2 Mo6S8 (100) surface. Using the NEB method, we calculate the activation energy to intercalate Mg from MgCl+ and plot the result in Figure 5. It is found that the energy cost to break the Mg−Cl bond on the CP (100) surface is only 0.2 eV. This is primarily because Mg2+ is significantly screened by the charge that accumulates on the neighboring S atoms that facilitates its intercalation,6 while Cl− is attracted by a surface Mo atom. Once the Mg−Cl bond is broken, the intercalation of the Mg ion resembles its bulk behavior in CP, with diffusion barriers of ∼0.5 eV, consistent with previous estimates.24,25 In addition, because of the relatively large ionic radius of the S atom, the three-dimensional open channels in the CP structure cannot co-intercalate MgCl+ as a unit. We note here that our estimated NEB barrier may be slightly different if the surface concentration of MgxCly+ changes. In this work, the surface concentration of MgxCly+ is limited to one unit per 168 Å2. For the case of Mg2Cl3+, we first test different adsorption geometries of Mg2Cl3+ on the CP (100) surface, as illustrated in Figure 6. It is found that the chainlike Mg2Cl3+ configuration (structures b and c) exhibits stronger surface adsorption because two Cl atoms are participating in surface Mo−Cl bonds. Compared to structure b, the Mg atoms in structure c are associated with more of the surface S atoms, resulting in an ∼0.4 eV lower energy state. Using structure c as the starting point, we perform a NEB simulation (as described above) and present the results in Figure 7. To break the first Mg−Cl bond (A to B) requires an activation energy of ∼0.7 eV, and a subsequent 0.8 eV activation barrier must be surmounted to

Figure 7. Calculated activation energy to intercalate Mg from a chainlike Mg2Cl3+ unit (structure c in Figure 6).

break the second Mg−Cl bond (B to C). In contrast with intercalation from MgCl+, where the highest energy barrier (slowest diffusion) is on the same order of magnitude as the bulk diffusion process, in Mg2Cl3+ the higher energy cost is related to the surface reactions. According to the Arrhenius equation, the ion diffusivity is proportional to exp(−Eact/kBT), where Eact is the activation energy. If Eact increases from 0.5 to 0.8 eV, Mg-ion diffusivity will decrease by almost 5 orders of magnitude at operating temperatures. This implies that the bottleneck of the entire Mg intercalation process may be related to the complex surface reactions rather than bulk properties of the cathode materials. Surface Contamination. From the NEB analysis described above, the stripping of Cl− relies on the competition between Mg and exposed surface Mo, i.e., which one is more willing to donate an electron to Cl. Once Mg is intercalated, Cl− will be left on the surface in the form of either Mo−Cl or MgCl2 5936

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Chemistry of Materials associated with two Mo sites, as shown in Figure 7. This will eventually passivate all Mo sites, which in turn could terminate further Mg intercalation. However, one possible mechanism that will allow continuous Mg intercalation is to have another electropositive element on the surface, such as Mg itself, to accommodate Cl−. For example, in Figure 7 (frame D), a neutral MgCl2 unit can be left on the surface and bind to two surface Mo atoms. This surface Mg, which is not intercalating, can interact with another incoming MgCl+ unit and restart the Mg intercalation process. To demonstrate this hypothesis, we optimize the surface structure with a (MgCl−MgCl2)+ cluster adsorbate, as shown in Figure 8, and simulate the Mg

Figure 9. Desorption energy of MgCl2 from various surface MgCly2−y species.

Figure 8. Calculated activation energy to intercalate Mg from a (MgCl−MgCl2)+ unit.

intercalation pathway using the NEB method. From Figure 8, the activation barrier to intercalate Mg is much higher than in previous cases. This is mainly because the surface Mg2+ is already strongly coordinated by two Cl− ions on the CP surface and thus less willing to adopt another Cl−. Although the surface Mg can also act as a cationic center on the CP surface, like Mo, to accept Cl− from the intercalated Mg2+, it takes more energy, ∼1.1 eV, to initiate the Mg intercalation reaction. From a bulk diffusion point of view, an activation barrier of 1.1 eV corresponds to a diffusion constant on the order of 10−20 cm2 s−1, which is far beyond the diffusion limit for practical battery applications. However, it is also possible that as the Cl-passivated CP surface continues to interact with the incoming MgCl+, it forms metastable MgCly2−y species on the surface, which increases the probability to release neutral MgCl2 species back into the electrolyte and frees up Mo sites on the CP surface. Here we compare the probability of releasing a neutral MgCl2 unit from a passivated CP surface by calculating the MgCl2 desorption energy from various surface configurations. An incoming MgCl+ can react with a Clpassivated CP surface by forming either a neutral MgCl2 or a negative MgCl3− species depending on the initial local Cl coverage, as shown in Figure 9. The energy cost to remove the neutral MgCl2 from the CP (100) surface is ∼1.7 eV because two strong Mo−Cl bonds have to be broken simultaneously. To release a MgCl2 unit from a charged MgCl3− species requires a lowered energy of 1.2 eV, and in this case, only one surface Mo−Cl bond is broken. Even though this energy of 1.2 eV is still quite high, we expect at real liquid−solid interfaces when the MgCl2 unit is better screened by the electrolyte and the solvent entropy contribution significantly increases, the energy cost to release MgCl2 back to solution can be greatly reduced.

Figure 10. Discharge profiles for CP cathode samples after 5, 10, 15, 20, and 25 cycles at 0.3 C.

These simple theoretical studies hint at the formation of complex MgxCly layers on the CP surface or the growth of large MgxCly oligomers (and subsequent precipitation) near the surface because of a diffusion-limited concentration gradient that overshoots the solubility limit for MgCl2 in THF. To confirm this hypothesis, Mg-ion full cells were constructed and cycled. In Figure 10, typical discharge curves are provided for the full cell experiment. After 5, 10, 15, 20, and 25 cycles, the cathode is removed and its surface is analyzed using image scanning electron microscopy with energy dispersive X-ray spectroscopy (SEM/EDX). The obtained post mortem SEM result for each cycled cell is shown in Figure 11. There is evidence of a foreign compound precipitating on the surface after as few as five cycles, which can be seen by the difference in contrast present in the secondary electron image. The darker objects contain lighter elements consistent with magnesium compounds rather than the bright background containing Mo. At five cycles, the compounds are very disperse and are visible in only the secondary electron image. SEM at 10 cycles shows a higher concentration of the compounds decorating the surface of the cathode; however, they are still very small in size. The SEM at 15 cycles shows clusters of strange rectangular crystallites that may have been formed via Ostwald ripening of the compounds present in the 5- and 10-cycle samples. To 5937

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concentration near the cathode surface would be enhanced by the reactions we propose. With little volume to carry the product away, local concentrations could quickly exceed the solubility limits, leading to precipitation at a much earlier time and leading to larger precipitates. We will continue to investigate this and other issues related to cell failure involving the Chevrel phase. Our work implies that in addition to the well-known stability issue of the electrolyte with respect to the anode material (Mg metal), which essentially prevents the use of aqueous solutions or organic solvent non-halide electrolytes with reasonable Coulombic efficiency, one should also consider the compatibility between the electrolyte and the cathode material. For example, with the hope of increasing the cell voltage, transition metal oxides have been proposed as the Mg intercalation cathode: such as thin-film V2O5, MoO3,3 and chemically delithiated spinel LiMn2O4.27 However, these insulating oxide structures usually do not present many open metal sites, but rather oxygen at their surfaces, which is unlikely to assist in dissociation of halide counterions from Mg2+ before intercalation through the cathode surface. At this point, maybe one should consider the following options for future development of Mg-ion battery cells: find transition metal compound cathode surfaces with more open metal sites, search for new metal cluster materials that are more likely to present the surface with unsaturated metal sites, and/or use less strongly coordinating anions in the electrolyte.

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CONCLUSION In this work, we study the surface properties of the Chevrel phase cathode, Mo6S8, and propose several Mg intercalation mechanisms through this cathode surface. The pristine CP (100) surface is chemically active and interacts with both the electrolyte salt (Mg x Cl y species) and solvent (THF) components. During cell discharge, electrons flow through the external circuit from the anode (Mg metal) into the cathode (CP), which makes the cathode material electron-rich. As a consequence, the CP surface is negatively charged and attracts positive ions, such as solvated MgCl+ and Mg2Cl3+, toward the surface. As these cations approach, the induced surface electron density of this metallic cathode increases. This strong localized negative charge on the CP surface will in turn weaken its interaction with the solvent molecules, because the oxygen lone pair of the ether group will be repelled from the negatively charged surface. This electronic screening effect predicted for CP will also benefit the Mg desolvation process. Using a cluster model, it is found that the Mg−Cl dissociation energy in a coordination complex involving THF is typically greater than 3 eV. However, the CP (100) surface acts as a catalyst and can reduce this dissociation energy to as low as ∼0.2 eV because Mg2+ is effectively screened by neighboring S atoms while Cl− binds to a surface Mo. Once all the open Mo sites on the surface are passivated by Cl−, further Mg intercalation exhibits higher intercalation barriers, ∼1 eV, apparently defining a surfacelimited process. However, over time, MgCl+ and Mg2Cl3+ species will continue to arrive at the cathode surface while waiting for Mg to intercalate, possibly forming layers of complex MgxCly species on the CP (100) surface. We have demonstrated in this work that these surface MgxCly layers do not necessarily prevent further Mg intercalation, and with some energy cost, neutral MgCl2 units can be released from the surface.

Figure 11. SEM (left) and corresponding secondary electron image (right) of the CP cathode in (a) 5, (b) 10, (c) 15, (d) 20, and (e) 25 cycles.

investigate the composition of these crystals, EDX was performed. Figure 12 displays the EDX elemental map of the crystallites in question. SEM confirms that at 15 cycles the foreign rectangular crystals are in fact some form of MgxCly. These crystals are rectangular in nature, which is consistent with the (110) plane of the MgCl2 crystal. The SEM analysis of the 20 cycle sample shows the beginning of three-dimensional growth of the rectangular crystallites while at 25 cycles the structures continue to grow and consume the cathode surface. One might reasonably ask why we observe the formation of these structures when other researchers have not seen similar results with cells of a similar makeup.26 As noted above, one key difference in our study is the use of a thinner polymeric separator. We hypothesize that this may have had an impact on the rate of formation of these structures as the cathode surface becomes coated with chloride. Because of the relatively small volume of electrolyte available, it may be that the local 5938

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Figure 12. Secondary electron image and the corresponding EDX map of the CP cathode cycled in APC for 15 cycles. Elements Mg, Cl, and Al are colored red, green, and blue, respectively. Al was mapped to serve as a background as it is present in the electrolyte but does not participate in precipitation.

supported by the Office of Science of the U.S. Department of Energy under Contract DE-AC02-05CH11231. The experimental work was done at Sandia National Laboratory, which is a multiprogram laboratory managed and operated by Sandia Corp., a wholly owned subsidiary of Lockheed Martin Co., for the U.S. DOE’s NNSA under Contract DE-AC04-94AL85000.

Depending on the starting concentration and the local diffusion constant of the solvated species, MgxCly oligomers may eventually precipitate. In reality, other surface geometries, much more complex than the (100) surface, also exist, especially for nanosized CP particles. Nevertheless, similar Mg intercalation behavior is expected because the surface will still be composed of Mo6 clusters surrounded by S atoms. As shown from our full cell tests, surface precipitates start to grow after as few as five cycles. This precipitation process is essentially irreversible and by the end of 25 cycles the cathode surface is completely covered by the MgxCly solid. This work has greater implications for multivalent cathode and electrolyte design in general. The requirement of halide electrolytes for compatibility with the Mg metal anode presents additional barriers to intercalation at the cathode that would persist in alternative higher-voltage transition metal oxides without the additional assistance of open metal sites at their surfaces or substitution of less coordinating anions in the electrolyte.

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ADDITIONAL NOTE The adsorption energy is estimated by comparing the total energies of an adsorbed MgxCly+CP structure with a structure in which the MgxCly adsorbate is extracted away from the surface to a distance of 7 Å in vacuum. The same strategy is used to compute the MgCl2 desorption energies reported in Figure 9. a

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REFERENCES

(1) Aurbach, D.; Lu, Z.; Schechter, A.; Gofer, Y.; Gizbar, H.; Turgeman, R.; Cohen, Y.; Moshkovich, M.; Levi, E. Prototype systems for rechargeable magnesium batteries. Nature 2000, 407, 724−727. (2) Yoo, H. D.; Shterenberg, I.; Gofer, Y.; Gershinsky, G.; Pour, N.; Aurbach, D. Mg rechargeable batteries: an on-going challenge. Energy Environ. Sci. 2013, 6, 2265−2279. (3) Gershinsky, G.; Yoo, H. D.; Gofer, Y.; Aurbach, D. Electrochemical and Spectroscopic Analysis of Mg2+ Intercalation into Thin Film Electrodes of Layered Oxides: V2O5 and MoO3. Langmuir 2013, 29, 10964−10972. (4) Wang, R. Y.; Wessells, C. D.; Huggins, R. A.; Cui, Y. Highly Reversible Open Framework Nanoscale Electrodes for Divalent Ion Batteries. Nano Lett. 2013, 13, 5748−5752. (5) Huie, M. M.; Bock, D. C.; Takeuchi, E. S.; Marschilok, A. C.; Takeuchi, K. J. Cathode materials for magnesium and magnesium-ion based batteries. Coord. Chem. Rev. 2015, 287, 15−27. (6) Thoele, F.; Wan, L. F.; Prendergast, D. Re-examining the Chevrel phase Mo6S8 cathode for Mg intercalation from an electronic structure perspective. Phys. Chem. Chem. Phys. 2015, inpress DOI: 10.1039/C5CP03046C. (7) Wan, L. F.; Prendergast, D. The Solvation Structure of Mg Ions in Dichloro Complex Solutions from First-Principles Molecular

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the Joint Center for Energy Storage Research, an Energy Innovation Hub funded by the U.S. Department of Energy (DOE), Office of Science, Basic Energy Sciences. The computational work was supported by a User Project at The Molecular Foundry using the computing cluster (vulcan), managed by the High Performance Computing Services Group, at Lawrence Berkeley National Laboratory (LBNL) and the resources of the National Energy Research Scientific Computing Center, LBNL, both of which are 5939

DOI: 10.1021/acs.chemmater.5b01907 Chem. Mater. 2015, 27, 5932−5940

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DOI: 10.1021/acs.chemmater.5b01907 Chem. Mater. 2015, 27, 5932−5940